One of the most important advances ever made in chemistry was the synthesis of ammonia from nitrogen and hydrogen gases. The process earned its discoverer (Fritz Haber) a Nobel Prize. It was important because up until that time, starvation had been a very real threat for most of the world, as a result of a population boom combined with failing harvests. Fritz Haber realized that the only real way to progress would be to make “bread out of air” – to turn the plentiful nitrogen in ordinary breathing air (78% nitrogen, by weight) into something they could use as fertilizer (ammonia, which would be oxidized to form the nitrates needed for plant growth). His process was a success, and mass starvation was avoided.
While his process was immediately commercialized and went on to make huge amounts of money for Haber and his industrial partners, it was still far from an ideal reaction. It required enormous temperatures and pressures, which pushed the steel industry at the time to their absolute limits in terms of what they could handle. Accidents and breakdowns were common. The stumbling block for Haber, and for scientists today who are trying to improve on his 100 year old process, is the strength of the chemical bonds that are holding the nitrogen molecule together. Nitrogen in the air exists in the form of two nitrogen atoms, bound together into a molecule. Since each ammonia molecule requires only one nitrogen, those two atoms must be broken apart before they can be used for synthesis. Unfortunately, the nitrogens are bound together with not one, not two, but three chemical bonds, making the separation of the two atoms incredibly difficult (hence the enormous energy required by the Haber process).
In the past year, there has been a fundamental laboratory advance, as reported by Cornell University in the journal Nature Chemistry. They have developed a method to split the nitrogen molecule into its two component parts, and the split is performed at room temperature and at normal pressure. The result is a process that is much less energy-intensive, and opens the possibility of using simple atmospheric nitrogen as the starting point for organic synthesis. Not only could it serve as the starting point for ammonia, but as the starting material for any nitrogen-containing organic compound. This approach is very exciting, and so far the only company who has managed to commercialize this process (nitrogen converted into a nitrogen-containing organic molecule) has been a tiny company in New York who set up shop next to Niagra Falls. It was only by directly harnessing the hydroelectric power from Niagra that they were able to generate enough electricity to power the reaction – it took an incredible force to break up the nitrogen molecule.
This discovery of a room temperature, normal pressure process is even more important when you consider that nitrogen-containing organics are common structures in most modern pharmaceuticals. Most medicines have at least one nitrogen incorporated somewhere in the molecule; providing a cheaper source of that fragment for the organic chemists to use would lower the cost of producing the material. Nitrogen-carbon bonds are also found in important consumer products like nylon, and insecticides. The Cornell researchers have developed a way of reacting carbon monoxide (a cheap, readily available gas found in the gas outputs of coal power plants) with nitrogen gas from the atmosphere to produce these important nitrogen-containing fragments.
Normally such a reaction also typically requires large amounts of energy. Chemists typically rely on a molecule having a negative or positive charge distribution, called a dipole, in order to find a point of attack for their synthesis. Nitrogen shares its electrons equally, which makes it exceptionally unreactive. The key to Cornells approach is a metal complex which contains the metal hafnium. In the first step of the reaction, the hafnium complexes surround a nitrogen molecule as sort of a cage. A reaction takes place between the hafnium and the nitrogen, which breaks two of the bonds and produces an intermediate. The carbon monoxide is then added, which reacts with the nitrogen-hafnium complex; upon addition of acid, the new nitrogen-containing molecule is freed from the “cage” of the hafnium and the final product is reached.
This reaction isn’t quite ready for commecial use just yet. The reaction isn’t catalytic, meaning that the hafnium complex gets used up during the reaction. If researchers can discover a way to recycle the hafnium, maybe by convincing the ligands to “let go” at the end of the reaction, the process will turn catalytic and will be ready for large-scale use. Other metals may also be possible, which would reduce the cost of the process as well. This is an exciting time for chemistry; we may soon see the process of Fritz Haber – nearly 100 years old at this point – finally retired.
Source: Nature Chemistry 2, 30 – 35 (2010) (see link, attached).